Question

Question 1: ​When creating the calibration curve during week 1, it was determined that [KSCN] = [FeSCN​2+​] at equilibrium. Explain why these two concentrations were equal in terms of the relative concentrations of each reactant, the limiting reactant, and the direction in which the reaction proceeded

Question 2: ​Describe the process by which equilibrium concentrations of the reactants and product were determined. Include in your discussion, how the equilibrium (final) concentration of FeSCN​2+​ was determined using a calibration curve, mole-to-mole ratios and how an ICE table should be completed. Explain which values were subtracted and added, and why

Question 3: ​What was the experimentally determined K​eq​ average? What is the theoretical K​eq​ for this reaction? How does your experimental value compare to the theoretical value and what are some factors that may have played a role in causing differing values

Question 5: ​Based on your experimentally determined K​eq​ value, are products or reactants favored? Explain your reasoning in terms of what it means to have a relatively small K​eq​ versus a relatively large K​eq​?

Answer 1

1. The equilibrium reaction is given as

Fe³⁺ (aq) + SCN^- (aq) ---------> FeSCN²⁺ (aq)

The equilibrium constant is given as

K_c = [FeSCN²⁺]/[Fe³⁺][SCN^-]

The concentration of Fe³⁺ is so chosen that [Fe³⁺] >> [SCN^-]. In order to keep K_c constant, [FeSCN²⁺] must increase to neutralize the effect of increased [Fe³⁺].

As per the stoichiometric equation, Fe³⁺ and SCN^- react on a 1:1 molar ratio. Again,

1 mole Fe³⁺ = 1 mole SCN^- = 1 mole FeSCN²⁺.

Since [Fe³⁺] >> [SCN^-], hence, the amount of FeSCN²⁺ formed in moles/L is equal to the molar concentration of SCN^-. The equilibrium reaction essentially proceeds in the forward direction giving more FeSCN²⁺.

2. In the first part of the experiment, the calibration curve is plotted by reacting excess Fe³⁺ with SCN^- to give FeSCN²⁺. A series of standard FeSCN²⁺ solutions are prepared by taking a large excess of Fe³⁺ and varying amounts of SCN^-. The absorbances of the standard solutions are measured and plotted against the known concentrations of SCN^- (which is equal to [FeSCN²⁺]) to produce a calibration curve with regression equation y = mx + c where m is the slope of the plot.

In the second part of the experiment, the absorbances of the unknown FeSCN²⁺ solutions are measured and fitted in the regression equation to get the concentration of unknown FeSCN²⁺.

Let us start with known concentrations x of Fe³⁺ and SCN^- and let the equilibrium concentration of FeSCN²⁺ be calculated as y. Therefore, the ICE chart is set up as

Fe³⁺ (aq) + SCN^- (aq) ---------> FeSCN²⁺ (aq)

initial                                 x                x

change                              -y               -y                                   +y

equilibrium                   (x – y)         (x – y)                                y

The value of the [FeSCN²⁺] was subtracted from the initial concentrations of Fe³⁺ and SCN^- because only a small portion of the reactants reacted to produce FeSCN²⁺ and the major portion was left as excess.

3 + 4. Need experimental data.